Chemistry: Properties and Application of Sb-Antimony
Abstract
Antimony is a semi metallic chemical element in Period 5 and Group 15 of the Periodic Table of chemical elements. Because it is semi metallic, it exists as both a metal and non-metal. The metallic form is silvery, bright, brittle and hard while the non-metallic form is a grayish powder. Like many elements in its group, antimony is a poor conductor of heat and electricity. Existence of antimony has been known since ancient times and was used by ancient alchemists to make other metals, beauty products, and concoctions for healing. Antimony is found in trace quantities in nature but is chiefly obtained from mineral ores such as valentinite (Sb2O3) and stibnite (Sb2S3). Pure refined antimony is used to make semiconductor devices such as infrared detectors and diodes. It is also alloyed with lead to make the latter more durable.
Introduction
Antimony is a shiny, silvery white element. Its surface is scaly, and is brittle and hard like a non-metal. As a metalloid, it exhibits characteristics of both metals and non-metals. Compounds of antimony have been used by human beings since ancient times. In ancient Egypt, women used the stibic stone (antimony sulfide) as make-up for eyes. The stibic stone was also used in making glassware and glazes for beads (Randich et al, 2002). The chemical name for antimony (stibium) was taken from the ancient Egyptian name for the element. Antimony is believed to have been named by Roman scholar Pliny (23–79 CE) who called it stibium. Arab alchemist Abu Musa Jabir Ibn Hayyan (721–815 CE) probably first called it antimony — ‘anti’ meaning not and ‘monos’ meaning alone because the element does not occur alone in nature (Shotyk, Krachler & Chen, 2006).
Although it was used for a long time, it was not until the 17th century that antimony was recognized as a chemical element. The first modern detailed information about antimony was published in the 18th century when French chemist Nicolas Lemery wrote the Treatise on Antimony (Krebs, 2006). Antimony occurs in two natural isotopes, antimony-121 and antimony-123. Isotopy occurs when two or more forms of an element differ from each other in terms of mass numbers. In addition, about twenty different radioactive isotopes of antimony are known. These isotopes give off some form of radiation. Two of these radioactive isotopes (antimony-124 and antimony-125) are used commercially as tracers (Emsley, 2011).
General Properties and Reactions
The chemical and physical properties of antimony are summarized in Table 1 below. Antimony metal is generally stable under normal conditions and does not react with air or water. It is also a poor conductor of electricity and heat. In the electrochemical series, antimony is positioned after hydrogen, meaning that it cannot displace ions of hydrogen from dilute acids. Simple cations of antimony (Sb+3 and Sb+5) do not occur in solutions but in hydrolyzed form such as Sb(OH)6¯ (Randich et al., 2002). The dominant cations species in the pH range unique to natural environments is Sb(OH)3 and Sb(OH)6¯ for trivalent antimony and pentavalent antimony respectively. In oxidizing environments, Sb(OH)3 is the dominant species under relatively reducing conditions whereas Sb(OH)6¯ tends to be dominant for pH values greater than 3 (Schmitt, 1960).
In natural water, the concentration of antimony has been found to be too low for antimony pentoxide (Sb2O5) or antimony trioxide (Sb2O3) to precipitate out. Antimony trioxide exhibits dimorphic properties, existing as an orthorhombic form (valentinite) or as a cubic form (senarmontite). The latter form is stable below 570°C of temperature. Additionally, antimony trioxide is amphoteric, meaning that it is soluble in hydrochloric acid, bases and certain organic acids but not in dilute nitric or sulphuric acid. Strong oxidizing agents such as nitric acid convert antimony trioxide into antimony pentoxide, which is strongly acidic (Schmitt, 1960).
Another major property of Antimony is that it forms complex ions with both organic and inorganic acids, the best known of which is tartrate. Stibine (SBH3) is one of the few gaseous antimony compounds. In this compound, the antimony is in the -3 valence state (Shotyk, Krachler & Chen, 2006). The compound is formed by the reactive effect of acids on antimony alloys or metal antimonides, electrolysis of basic or acidic solutions where antimony is used as a cathode, or reduction of antimony compounds. This implies that there is the danger of stibine being precipitated from lead storage batteries where antimony is alloyed with the lead. With time, stibine decomposes into hydrogen and metallic antimony. It is readily oxidized by air at normal conditions to form water and antimony trioxide (Krebs, 2006).
Electrolytic deposition of antimony results in an unstable, amorphous form of the element called explosive antimony. When scratched or bent, explosive antimony changes mildly in a very explosive manner to the stable metallic form. There is also a yellow form of antimony that results from mild temperature oxidation of stibine, and an amorphous black that results from sudden quenching of the vapor. Metallic antimony does not react with moisture or air under ordinary conditions, but is readily converted into oxide if the air is moist. The halogens and sulphur can easily oxidize antimony when heated (Emsley, 2011).
The electronic structure of antimony is closely related to that of arsenic, and consists of three half-filled orbitals in the last shell. It is thus able to form covalent bonding and exhibits -3 and +3 oxidation states (Haynes, 2015). Antimony acts as an oxidizing agent and readily reacts with many metals to form antimonides. All antimonides, in general, resemble phosphides, nitrides and arsenides but are somehow more metallic. For the purpose of analytical chemistry, antimony can be easily weighed and separated for analysis as the antimony sulfide, (Sb2S3). In an alternative process, the sulfide is converted to oxide and then weighed as Sb4O6. In addition, there is a wide range of volumetric methods such as oxidization of antimony with potassium permanganate iodine or potassium bromated. The modified Gutzeit method can be used to determine small amounts of antimony (Harder, 2002).
Summary Physical Properties of Antimony
Chemical symbol
Sb
Atomic Number
51
Atomic weight
121.760
Melting point
903.78K ((630.63°C or 1167.13°F))
Boiling point
1860 K (1587°C or 2889°F)
State at room temperature
Solid
Elemental classification
Semi-metal
Period
5
Group
15
Group name
Pnictogen
Density
6.684 g/cm3
Ionization Energy
8.64 eV
Oxidation States
+5, +3, -3
Electron configuration
1s22s22p63s23p63d104s24p64d105s25p3
Occurrence
Antimony is hardly found in its native state (as an element). Instead, it occurs as a compound in more than 100 different minerals. The most common minerals containing antimony are stibnite, bournonite, tetrahedrite, jamesonite and boulangerite. In most of these minerals, antimony is found combined with sulfur to form antimony sulfide (Sb2S3). Other major commercial minerals of antimony are cervantite, stibcontite, kermasite, valentinite and senarmontite. Complex ores such as livingstonite are also a major source of antimony. The abundance of antimony in the earth’s crust is estimated to be 0.2 parts per million, making it one of the rarest of the chemical elements found in the earth’s crust. China, Russia, Kyrgyzstan and South Africa are the largest producers of antimony in the world. The United States produces substantial quantities of antimony as a by-product at the Idaho silver mine (Haynes, 2015).
Uses and Applications of Antimony
Antimony is mainly used metallurgically as an additive substance because its physical properties are not suitable for engineering. By far, its most important commercial use is as an alloying component for lead and certain lead-based alloys to improve corrosion resistance and make the alloy hard and stiff. Antimony is also used as an alloying constituent in tin to produce tin-based babbits and pewter for use in bearing metal applications. The element is also widely used in the manufacture of castings, soldiering materials and cable insulations. Certain lead-antimony alloys are used in the manufacture of low friction metals, batteries and type metals among other commercial products. Other antimony compounds are used to manufacture paints, flame proofing materials, glass, ceramic enamels and pottery (Harder, 2002).
Major Applications with all Structural and Functional Details
The structural properties of antimony and its compounds make it suitable for use in a variety of other industrial and commercial applications. The most common compound, antimony sulfide, is used to vulcanize rubber. Its unique chemical properties make it ideal for use as a vermillion pigment and certain other pigment shades such as orange and yellow, which are formed by slow oxidation of the sulfide. To a lesser extent, antimony sulfide is used in fireworks, racer bullets and ammunition primers. Pure antimony (purity level exceeding 99.999%) is applied in semi-conductor technology. Such high levels of purity can be obtained from the reduction of high purity compounds such as chloride and trioxide with hydrogen. Important compounds of antimony with groups III or V oxidation state as (AlSb, GaSb and InSb) are widely used as diodes, infra red detectors, and Hall effect devices (Robert, 2006).
Antimony and its compounds are also used in the field of medical sciences. Antimony trioxide is used in the preparation of certain medicines called antimonials, which are used as emetics. Selected antimony compounds are used in the treatment of protozoans. Tarter emetic (Potassium antimonyl tartrate) was once used as a leading anti-schistosomal drug but has been replaced by praziquantel. Antimony and some of its compounds are used to prepare veterinary medicines such as lithium antimony thiomalate, which is applied to ruminants as a skin conditioner. In other animals, antimony is used because of its keratinized tissues. Antimony can be toxic depending on its chemical state. Generally, metallic antimony is inert but stibunite is very toxic. When handling antimony and its compounds, proper ventilation should be used to avoid contamination. Notable cases of dermatitis and other skin conditions have been reported in facilities handling antimony (Haynes, 2015).
Oxidation States Exhibited by Antimony
Shells
2,8,18,18,5
Electron configuration
[Kr] 4d10 5s2 5p3
Minimum oxidation number
3
Maximum oxidation number
5
Minimum common oxidation number
0
Maximum common oxidation number
5
Electronegativity (Pauling Scale)
2.05
Polarizability volume
6.6 Å3
Structure and Coordinate Geometry
Antimony’s coordinate geometry consists of three covalent bonds and a single lone pair of electrons. Reluctance of this element to engage in hybridization results in typical bond angles approximating 90 degrees. Subsequent steric interactions and chelated structures have been shown to enforce other geometries. The tetracoordinate tetrahedral geometry shown in Figure 1 below is common for antimony and other elements. Computational studies of antimony’s potential bonding arrangements suggest minimal d-orbital participation. Thus, antimony’s double bond arrangement is described as singly bonded with localized negative charges (Harder, 2002).
Figure 1: Bonding Model For Antimony Coordinate Geometry
Conclusion
Antimony and its mineral compounds have been known since ancient times. Antimony is one of the numerous elements that occur naturally in the environment, although it is found mainly in compound form. It also finds its way into the environment through numerous applications by human actions. Antimony was discovered as an element in the 17th century although it had been in use several centuries earlier. Ancient Egyptians and Roman alchemists used antimony to prepare beauty products and drugs. In the world economy, antimony is an important commercial element. It is used in the manufacture of various industrial products and as an alloy for strengthening other metals such as lead and tin. Russia and China are the leading producers of antimony.
Being a rare element, antimony is mainly found in its natural form of sulfide stibnite. A highly pure form of antimony is used to make different types of semiconductor devices such as infrared detectors and diodes. Alloys of antimony and lead are used to manufacture batteries, cable sheatings, bullets and other products such as glass and paints. Antimony is also widely used in making flame-retardant materials. About half of all antimony alloys goes to this use. Both antimony and its compounds are highly toxic and thus dangerous to human health. Even in low levels, it can cause an irritating sensation to the lungs and eyes. It can also cause stomach pain, vomiting and ulcers. At higher doses, antimony contamination can cause severe organ failure and even death.
References
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Haynes, W. (2015). CRC Handbook of Chemistry and Physics. CRC Press/Taylor and Francis, Boca Raton.
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Schmitt, H. (1960). Determination of the energy of antimony-beryllium photoneutrons. Nuclear Physics 20: 220.
Shotyk, W., Krachler, M. & Chen, B. (2006). Contamination of Canadian and European bottled waters with antimony from PET containers. Journal of Environmental Monitoring 8 (2): 288–92.
